Boron — Important Compounds (Borax, Boric Acid, Diborane)

Why Boron's Compounds Are Different

Boron heads Group 13 with the valence configuration $\ce{[He] 2s^2 2p^1}$ — only three electrons available for bonding, distributed over four valence orbitals ($2s$, $2p_x$, $2p_y$, $2p_z$). The old-NCERT supplement traces almost every peculiarity of its compounds to this single fact: the trivalent state leaves only six electrons around boron, two short of an octet. Such electron-deficient centres are hungry for an electron pair, so boron compounds behave as Lewis acids, accepting a lone pair to complete the octet.

A second restriction follows from boron being a second-period element with no accessible $d$ orbitals: its maximum covalence is fixed at four. Boron can form $\ce{[BF4]^-}$ but never $\ce{[BF6]^3-}$, whereas aluminium, which has $3d$ orbitals available, gives $\ce{[AlF6]^3-}$. Keep these two themes — electron deficiency and a covalence ceiling of four — in mind; they explain the structure and reactions of all three compounds below.

Borax — Structure and Bead Test

Borax is the most important compound of boron — a white crystalline solid usually written as $\ce{Na2B4O7.10H2O}$. That common formula hides the real structure. Borax actually contains the tetranuclear anion $\ce{[B4O5(OH)4]^2-}$, so its correct formula is $\ce{Na2[B4O5(OH)4].8H2O}$. Within this anion, two boron atoms sit in trigonal $\ce{BO3}$ ($sp^2$) units and two sit in tetrahedral $\ce{BO4}$ ($sp^3$) units — but in every form there are four boron atoms per formula unit.

Borax dissolves in water to give an alkaline solution, because the borate hydrolyses to sodium hydroxide and the weak orthoboric acid:

$$\ce{Na2B4O7 + 7H2O -> 2NaOH + 4H3BO3}$$

The borax bead test

On heating, borax first loses its water of crystallisation and swells up. On stronger heating it melts to a transparent liquid that solidifies into a glassy borax bead of sodium metaborate and boric anhydride:

$$\ce{Na2B4O7.10H2O ->[\Delta] Na2B4O7 ->[\Delta] 2NaBO2 + B2O3}$$

The metaborates of many transition metals carry characteristic colours, so the bead becomes a quick laboratory test. Heating borax on a platinum-wire loop with $\ce{CoO}$ in a Bunsen flame, for instance, gives a blue $\ce{Co(BO2)2}$ bead.

Orthoboric Acid — A Lewis Acid

Orthoboric acid, $\ce{H3BO3}$ (equivalently $\ce{B(OH)3}$), is a white crystalline solid with a soapy touch. It is sparingly soluble in cold water but dissolves freely in hot water. It is prepared by acidifying borax — with sulphuric acid in the NIOS route, or with hydrochloric acid in the NCERT route:

$$\ce{Na2B4O7 + H2SO4 + 5H2O -> Na2SO4 + 4H3BO3}$$

$$\ce{Na2B4O7 + 2HCl + 5H2O -> 2NaCl + 4B(OH)3}$$

The defining NEET point is its acid behaviour. Boric acid is a weak monobasic acid, but it is not a protonic (Brønsted) acid — it does not release $\ce{H+}$ from its own $\ce{O-H}$ bonds. Instead, the electron-deficient boron accepts a lone pair from a hydroxyl ion of water to complete its octet, and that abstraction of $\ce{OH-}$ is what leaves a hydronium ion behind:

$$\ce{B(OH)3 + 2H2O -> [B(OH)4]^- + H3O+}$$

The solid is built as a layered structure: planar $\ce{BO3}$ units in which each boron sits at the centre of three oxygens are linked into two-dimensional sheets through hydrogen bonds. The sheets stack with weak van der Waals forces between them, which is why the crystals cleave into flakes and feel soapy. Note that the hydrogen bonding here is intermolecular and gives a sheet, not a discrete molecule.

On heating, boric acid loses water in stages — first to metaboric acid near 370 K, then to boric oxide:

$$\ce{H3BO3 ->[\Delta] HBO2 ->[\Delta] B2O3}$$

Diborane and the Banana Bonds

Diborane, $\ce{B2H6}$, is the simplest known boron hydride and the most heavily examined structure in this topic. It is a colourless, highly toxic gas with a boiling point near 180 K that catches fire spontaneously on exposure to air.

Preparation

The laboratory and industrial routes from the two sources are tabulated below. Note that the NCERT supplement starts from $\ce{BF3}$ while the NIOS chapter also gives the $\ce{BCl3}$ route — both are valid.

Key reactions

Diborane burns in oxygen with a very large release of energy, and is readily hydrolysed by water back to boric acid:

$$\ce{B2H6 + 3O2 -> B2O3 + 3H2O}\qquad \Delta_c H = -1976~\text{kJ mol}^{-1}$$

$$\ce{B2H6 + 6H2O -> 2B(OH)3 + 6H2}$$

With Lewis bases it undergoes cleavage to borane adducts $\ce{BH3.L}$ — for example $\ce{B2H6 + 2NMe3 -> 2BH3.NMe3}$ and $\ce{B2H6 + 2CO -> 2BH3.CO}$. With ammonia and subsequent heating it gives borazine $\ce{B3N3H6}$, the so-called inorganic benzene.

Structure: the 3c-2e bridge bonds

Here is the crux. Diborane has eight $\ce{B-H}$ linkages but only twelve valence electrons — too few to make eight ordinary two-electron bonds. The structure resolves this as follows: the four terminal $\ce{B-H}$ bonds are normal two-centre two-electron (2c-2e) bonds, while the two bridging $\ce{B-H-B}$ linkages are three-centre two-electron (3c-2e) bonds — a single electron pair shared over three nuclei. Their curved geometry earns them the name banana bonds. The four terminal hydrogens and the two borons lie in one plane; the two bridging hydrogens sit symmetrically above and below it. Each boron is $sp^3$ hybridised.

From diborane, boron also forms hydridoborates; the most important is the tetrahedral $\ce{[BH4]^-}$ ion. $\ce{LiBH4}$ and $\ce{NaBH4}$ are made by reacting metal hydrides with $\ce{B2H6}$ in ether and serve as reducing agents in organic synthesis.

The Three Compounds Compared

A consolidated table is the fastest revision tool for this subtopic. Note how a single thread — boron's electron deficiency — runs through the alkaline borax solution, the Lewis acidity of boric acid, and the bridge bonding of diborane.